Why are single bonds weaker than double?

Question

Jacob Anderson · Accepted Answer

Simply from how they're constructed. Since a pure double bond consists of 1 #sigma# bond and 1 #pi# bond, it is one #\mathbf(pi)# bond's worth stronger than a single bond.

All pure single bonds consist of one #sigma# bond, i.e. due to one head-on orbital overlap.

Below is an example a constructive #2p_z-2p_z# head-on overlap that forms a #sigma_(2p_z)# molecular orbital, where electron density lies in the white bulged region---between the atoms.

A nice example is the #\mathbf(sigma)# bond in #\mathbf("Cl"-"Cl")#.
All pure double bonds consist of an additional #pi# bond, i.e. due to a sidelong orbital overlap.

Below is an example of the #2p_x-2p_x# constructive/bonding overlap, where electron density lies in the white bulged region (above the atoms).

An explicit #pi# bond example is the #\mathbf(pi)# bond in #\mathbf("O"="O")#, a product of either a #2p_x-2p_x# sidelong overlap, or a #2p_y-2p_y# sidelong overlap (but not both), that forms a #pi_(2p_(x"/"y))# orbital, depending on which pair overlaps.

This #pi# bond is made in addition to the #sigma# bond that was already made upon forming the first #"O"-"O"# bond.
Therefore, since a pure double bond consists of 1 #sigma# bond and 1 #pi# bond, it is one #\mathbf(pi)# bond's worth stronger than a single bond.

Paisley Boney · Answer

undefined Single bonds are weaker than double bonds because they consist of only one pair of electrons shared between two atoms, whereas double bonds consist of two pairs of electrons. The additional electron pair in a double bond provides greater electron density between the bonded atoms, leading to stronger electrostatic forces of attraction and a shorter bond length. As a result, double bonds require more energy to break compared to single bonds, making them stronger. Additionally, the presence of multiple bonds allows for greater overlap of atomic orbitals, which further strengthens the bond.